8-ATOMIC EMISSION SPECTRA

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1 Chemistry ATOMIC EMISSION SPECTRA Knowledge of the arrangement of electrons around the nuclei of atoms has been obtained by examining the light emitted by excited atoms. Atoms become excited when they absorb energy; they then emit energy in the form of light as they return to a less excited state. Under some conditions, the light emitted may be visible. When emitted light is reflected from a diffraction grating, only a few bright lines can be observed. These lines make up the atomic emission spectrum of the atom being studied. Each individual line represents a definite wavelength of light emitted by the excited atom. The diffraction grating is a device which separates light into a spectrum of its individual wavelengths. A prism behaves similarly. Gratings are made by etching fine, parallel, equallyspaced grooves on a glass plate. The diffraction gratings we will use are important parts of the spectroscopes, to be described below. Light is electromagnetic radiation and can be described in terms of waves with characteristic wavelengths. The relation between the wavelength (λ), speed (c), and frequency (υ) is given by the equation: λ υ = c (1) where c is the speed of light in a vacuum, x 10 8 m/s. The wavelength range for visible light is about 400 nm (violet) to 780 nm (red) corresponding to frequencies of about s 1 (1 nm = 10 9 m). In this experiment, we will examine the light emitted when hydrogen and helium absorb energy from an outside source and then emit energy in the form of light. The two elements are contained in glass tubes and a high-voltage power supply creates an electric discharge through the gas in the tube. In the case of hydrogen, the electric discharge dissociates the H 2 molecules into H atoms and excites the hydrogen electron. According to Bohr's theory of the hydrogen atom and quantum mechanical theory, there exists around the nucleus a series of energy levels in which the electrons may reside. In the ground (unexcited) state of an atom, the electrons occupy the lowest energy levels. Some electrons may be excited to higher levels by the absorption of specific amounts of energy. As the electrons return to their ground states, the absorbed energy is released, and if the sample contains many excited atoms, all possible energies of light for the atom will be emitted. The possible energies can be calculated from the difference of two electronic energy levels, and the relationship between the energy and the frequency and wavelength is given by the following equation: E final E initial = E = hυ = hc/λ (2) The quantity h is known as Planck's constant and is equal to x J s (J = joule, s = second). The emission spectrum consists of discrete lines corresponding to the differences in energy levels characteristic of, and unique to, the atoms of the element. The lines are normally tabulated by wavelength.

2 Bohr found that, for a one-electron atom or ion, the energy of an electron occupying an orbit with principal quantum number n is proportional to 1/n 2. For hydrogen, this relation is expressed by the equation: E n = -k(1/n 2 ) (3) where k for the hydrogen atom is the proportionality constant known as the Rydberg constant and has the value x J; n is the principal quantum number and can be any integer from 1 to infinity. When an electron drops from n initial to n final, the energy change is given by the expression: E = E final E initial = [ -k(1/n 2 ) final ] [ -k(1/n 2 ) initial ] = k [ (1/n 2 ) initial (1/n 2 ) final ] (4) From equation (2) we can calculate the wavelength or frequency corresponding to any energy difference. The values for Planck's constant and the Rydberg constant given above lead to a wavelength in meters. For convenience, we usually express the wavelength in nanometers. Figure 7 shows some of the electronic transitions possible for a hydrogen atom, and the type of light associated with the transition. Since n can take any value from 1 to infinity, there is an infinite number of transitions possible, but since the energy levels are closer together as n increases, there is a practical limit to the number of lines observed. Principal Quantum Number n=6 n=5 n=4 E N E R G Y n=3 n=2 Balmer Series (Visible and ultraviolet) Paschen Series (Infrared) n=1 Lyman Series (Ultraviolet) Figure 7: Approximate Electron Transition Energy-level Diagram for the Hydrogen Atom

3 THE SPECTROSCOPE The light source and spectroscope are shown schematically in Figure 8. The leads from the hydrogen or helium tube are connected to the terminals of a power source. There is a potential difference of 5000 volts between the terminals. Keep your hands away from the lamp and power source. Work only when the instructor is present. Light emitted from the tube enters the spectroscope through a narrow adjustable slit. By using the narrowest slit, a fine beam of light is examined. The beam is focussed on a diffraction grating which separates it into its component wavelengths and the resulting visible, colored lines are superimposed on a fixed scale calibrated to read in any convenient units. Our spectroscopes read in cm, that is, 4.25 on the scale means 4.25 x 10 5 cm. The conversion factors 1 cm = 10 2 m and 1 nm = 10 9 m can be used to change our reading to nm: 4.25 x 10 5 cm = 4.25 x 10 2 nm. Scale 5000 v Light beam Grating Port to admit light Eye PROCEDURE Figure 8: Measurement of Emission Spectrum Using a Spectroscope 1. Determine the wavelengths of the helium lines by carefully observing their positions on the scale. To avoid parallax error, be sure to look at each line in turn directly while reading the scale for that line. The scale will be brighter if you open the port on the left-hand side of the spectroscope. Record all of the wavelengths that are visible to you in your lab notebook. Some lines are very faint. Some of the wavelengths for helium are given in the table below. Some wavelengths for visible lines in the helium spectrum Color Wavelength in nm Red Yellow Green Blue-violet 447.1

4 2. Determine the wavelengths of the three readily visible hydrogen lines. A fourth line may be barely visible, depending on your own eye; record its wavelength if you see it. Compare the wavelengths you observe with the wavelengths calculated for the Balmer Series from equations (4) and (2). See Table 2 in the Pre-Laboratory Assignment. 3. Using the other setup in the laboratory with a diffraction grating and emission tube, and working with a partner, calculate the wavelengths of spectral lines from meter-stick measurements. The number of lines to be calculated will depend on the element being studied; consult the instructor. The setup is diagrammed in Figure 9. Light source x Observed line a θ L Grating Eye Figure 9: Line Observed with Diffraction Grating DO NOT LOOK DIRECTLY AT THE LIGHT SOURCE. Use a meter stick to measure distances a and x in Figure 3. x is the distance from the emission tube (the source) to a line in the spectrum. One partner, the observer, looks through the grating in order to see a line, and the other partner measures the distance to the point on the meter stick where his partner sees the line superimposed. The value of L can be calculated from: L 2 = a 2 + x 2 The angle θ is related to the wavelength (λ) of the line by x/l = λ/d = sin θ where d is the spacing between the lines on the diffraction grating. The spacings are usually given in lines per inch. The grating in our laboratory has lines/inch. Calculate sin θ from x and L; calculate d in cm/line. Then: λ = d sin θ

5 Chemistry ATOMIC EMISSION SPECTRA Section Report Sheet Name 1. Helium Spectrum (you may not actually see nine lines) Wavelength (nm) Line Color Experimental Corresponding Value Value from Table Hydrogen Spectrum Wavelength (nm) Line Color Experimental Value from Pre-Lab Value Diffraction Grating: Line measurement and calculation Line Element Number Color Wavelength Attach a separate sheet with your calculations.

6 Chemistry ATOMIC EMISSION SPECTRA Questions 1. If the lowest energy emission in the Lyman Series could be observed, would it lie in the visible range? Calculate its wavelength. 2. The Paschen Series of lines arises from electronic transitions to energy level n=3. Calculate the frequency in Hz of the light emitted by an electron falling from n=4 to n=3. 3. Explain why the emissions of the Paschen Series are lower energy overall than those in the Balmer Series.

7 Chemistry ATOMIC EMISSION SPECTRA Section Name Pre-Laboratory Assignment 1. Calculate the wavelengths of the lines in the Balmer Series from the values of n initial and n final using equations (4) and (2). Give the wavelength in nanometers. Show a sample calculation here, and enter the values in the table below. Electronic Transitions in the Spectrum of Hydrogen Line n i n f 1/n i 2-1/n f 2 Wavelength, nm a. For hydrogen, ionization would correspond to moving the electron from n=1 to. Using equation (4), calculate the ionization energy of one atom of hydrogen in joules. b. Calculate the ionization energy of atomic hydrogen in kj/mol.